David A. Katz Chemist, Educator, Science Communicator, and Consultant Department of Chemistry, Pima Community College

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1 WRITING CHEMICAL EQUATIONS 2004, 2002, 1989 by David A. Katz. All rights reserved. Permission for classroom used provided original copyright is included. David A. Katz Chemist, Educator, Science Communicator, and Consultant Department of Chemistry, Pima Community College I. THE MEANING OF A CHEMICAL EQUATION A chemical equation is a chemist s shorthand expression for describing a chemical change. As an example, - consider what takes place when iron rusts. The equation for this change is: Fe + O 2 Fe 2 O 3 In this expression, the symbols and formulas of the reacting substances, called the reactants, are written on the left side of the arrow and the products of the reaction are written on the right side. The arrow is read as gives, yields, or forms and the plus (+) sign is read as and. When the plus (+) sign appears between the formulas for two reactants, it can be read as reacts with. (The + sign does not imply mathematical addition.) The equation, above, can be read as iron reacts with oxygen to yield (or form) iron(iii) oxide. II. BALANCING A CHEMICAL EQUATION As it is written, the equation indicates in a qualitative way what substances are consumed in the reaction and what new substances are formed. In order to have quantitative information about the reaction, the equation must be balanced so that it conforms to the Law of Conservation of Matter. That is, there must be the same number of atoms of each element on the right hand side of the equation as there are on the left hand side. If the number of atoms of each element in the equation above are counted, it is observed that there are 1 atom of Fe and 2 atoms of O on the left side and 2 atoms Fe and 3 atoms of O on the right. Fe + O 2 Fe 2 O 3 Left side: Right side: 1 atom Fe 2 atoms Fe 2 atoms O 3 atoms O The balancing of the equation is accomplished by introducing the proper number or coefficient before each formula. To balance the number of O atoms, write a 3 in from of the O 2 and a 2 in front of the Fe 2 O 3 : Fe + 3 O 2 2 Fe 2 O 3 The equation, above, now has 6 atoms of O on each side, but the Fe atoms are not balanced. Since there is 1 atom of Fe on the left and 4 atoms of Fe on the right, the Fe atoms can be balanced by writing a 4 in front of the Fe: 4 Fe + 3 O 2 2 Fe 2 O 3 This equation is now balanced. It contains 4 atoms of Fe and 6 atoms of O on each side of the equation. The equation is interpreted to mean that 4 atoms of Fe will reaction with 3 molecules of O 2 to form 2 molecules of Fe 2 O 3.

2 It is important to note that the balancing of an equation is accomplished by placing numbers in front of the proper atoms or molecules and not as subscripts. In an equation, all chemical species appear as correct formula units. The addition (or change) of a subscript changes the meaning of the formula unit and of the equation. Coefficients in front of a formula unit multiply that entire formula unit. Another example of balancing an equation is: Al(OH) 3 + H 2 SO 4 Al 2 (SO 4 ) 3 + H 2 O Counting the atoms of each element in the equation it is found that there are 1 atom Al, 7 atoms O, 5 atoms H, and 1 atom S on the left side and 2 atoms Al, 13 atoms O, 2 atoms H, and 3 atoms S on the right side. Al(OH) 3 + H 2 SO 4 Al 2 (SO 4 ) 3 + H 2 O Left side: Right side: 1 atom Al 2 atoms Al 7 atoms O 13 atoms O 5 atoms H 2 atoms H 1 atom S 3 atoms S The counting, however, can be simplified by observing that the S and O in the SO 4 polyatomic ion acts as a single unbreakable unit in this equation. Recounting, using the SO 4 as a single unit, it is found that there are 1 atom Al, 3 atoms O, 5 atoms H, and 1 SO 4 polyatomic ion on the left side and 2 atoms Al, 1 O atom, 2 H atoms, and 3 SO 4 polyatomic ions on the right side. Al(OH) 3 + H 2 SO 4 Al 2 (SO 4 ) 3 + H 2 O Left side: Right side: 1 atom Al 2 atoms Al 3 atoms O 1 atoms O 5 atoms H 2 atoms H 1 SO 4 group 3 SO 4 groups Starting with Al, the atoms of Al can be balanced by writing a 2 in front of the Al(OH) 3 : 2 Al(OH) 3 + H 2 SO 4 Al 2 (SO 4 ) 3 + H 2 O Looking at the SO 4 ions, these are balanced by writing a 3 in front of the H 2 SO 4 : 2 Al(OH) H 2 SO 4 Al 2 (SO 4 ) 3 + H 2 O Now, only the O atoms and H atoms remain unbalanced. There are 6 atoms of O and 12 atoms of H on the left hand side of the equation and only 1 atom O and 2 atoms H on the right side. These can be balanced by writing a 6 in front of the H 2 O: 2 Al(OH) H 2 SO 4 Al 2 (SO 4 ) H 2 O The equation is now balanced and it is interpreted to mean that 2 molecules of Al(OH) 2 react with 3 molecules of H 2 SO 4 to form 1 molecule of Al 2 (SO 4 ) 3 and 6 molecules H 2 O. 2

3 Problems: Balancing chemical equations Balance each of the following equations: 1. H 2 + Br 2 HBr 2. N 2 + H 2 NH 3 3. Sb + O 2 Sb 4 O 6 4. Cu(NO 3 ) 2 CuO + NO 2 + O 2 5. (NH 4 ) 2 Cr 2 O 7 Cr 2 O 3 + N 2 + H 2 O 6. C 2 H 6 + O 2 CO 2 + H 2 O 7. Al + HgCl 2 AlCl 3 + Hg 8. FeS + O 2 Fe 2 O 3 + SO 2 9. KOH + Cl 2 KCl + KClO + H 2 O 10. Ca(OH) 2 + H 3 PO 4 Ca 3 (PO 4 ) 2 + H 2 O 11. BaCl 2 + Na 2 SO 4 BaSO 4 + NaCl 12. CrBr 3 + Na 2 SiO 3 Cr 2 (SiO 3 ) 3 + NaBr Not all equations can be easily balanced by the method used here. In some equations the oxidation numbers of some atoms change during the reaction. Such equations are known as oxidation-reduction equations and many of these require special methods to balance them. Although the balancing of oxidation-reduction equations will not be covered in this tutorial, the following oxidation reduction equation is provided as an exercise: Balance the following equation: Cu + HNO 3 Cu(NO 3 ) 2 + NO + H 2 O 3

4 III. TYPES OF CHEMICAL REACTIONS Most inorganic reactions can be classified into one of five general categories: direct union or combination, decomposition, displacement, metathesis or double displacement, and combustion reactions. Each of these will be discussed in more detail in the following sections. 1. Direct Union or Combination Reactions Any reaction in which two or more substances combine to form a single product is a direct union or combination reaction. The general form of a direct union reaction is A + B AB This type of reaction generally takes place between the following types of compounds: a. A metal + non-metal b. Metal oxide + non-metal oxide 2 Na + Cl 2 2 NaCl sodium chloride Fe + S FeS iron(ii) sulfide K 2 O + SO 3 K 2 SO 4 potassium sulfur potassium oxide trioxide sulfate c. Non-metal + non-metal CaO + CO 2 CaCO 3 calcium carbon calcium oxide dioxide carbonate C + O 2 CO 2 carbon dioxide N Cl 2 2 NCl 3 nitrogen trichloride 2. Decomposition Reactions Decomposition is the reverse of combination. That is, a single reactant is broken down into two or more products either elements or compounds. A decomposition reaction will take place because the compound is unstable or as a result of heating or electrical decomposition (electrolysis). The general form for a decomposition reaction is: AB A + B Some examples of decomposition reactions are: 4

5 2 HgO 2 Hg + O 2 mercury(ii) oxide 2 KClO 3 2 KCl + 3 O 2 potassium potassium chlorate chloride CaCO 3 CaO + CO 2 calcium calcium carbon carbonate oxide dioxide To understand how to predict products of decomposition reactions, see Section V. The Effect of Heat on Metallic Compounds, page Displacement Reactions (Sometimes called oxidation-reduction equations) A displacement reaction involves an element reacting with a compound whereby the element displaces a second element from the compound. The general form of this type reaction is: A + BC AC + B Displacement reactions usually occur between the following combinations: a. An active metal + an acid When a metal which is above hydrogen in the activity series is reacted with an acid, hydrogen is liberated and a salt is formed. (Refer to Section IV, The Electromotive Series, page 9) Zn + 2 HCl ZnCl 2 + H 2 hydrochloric zinc acid chloride b. A metal + a salt Mg + H 2 SO 4 MgSO 4 + H 2 sulfuric magnesium acid sulfate Each metal in the activity series displaces any metals below it to form a salt in solution. (Refer to Section IV, The Electromotive Series, page 9) Cu + 2 AgNO 3 silver nitrate Cu(NO 3 ) Ag copper(ii) nitrate Fe + CuSO 4 FeSO 4 + Cu copper(ii) iron(ii) sulfate sulfate 5

6 c. A Halogen + halide salt A halogen (F, Cl, Br, I, At) will displace any less active halogen from a halide salt. The order of activity decreases going from top to bottom down the halogen family in the periodic table. Cl NaI 2 NaCl + I 2 sodium sodium iodide chloride 4. Metathesis or Double Displacement Reactions A metathesis is a double displacement reaction that usually occurs in water solution. The general form of a metathesis reaction is: AB + CD AD + CB In order to have any appreciable degree of completion of metathesis reactions, one or both of the products must become unavailable for the reverse reaction. the principal conditions that favor the completion of these reactions are: (1) Formation of an insoluble compound - a precipitate (2) Formation of a gas (3) Formation of water Metathesis reactions are generally classified as precipitation reactions or as neutralization reactions. a. Precipitation Reactions In this type of reaction, two compounds which are water soluble react to form two new compounds, one of which is a precipitate (i.e. insoluble in water). The precipitate is often indicated by an arrow pointing downward,, written next to its formula. AgNO 3 + NaCl AgCl + NaNO 3 silver sodium silver sodium nitrate chloride chloride nitrate BaCl 2 + K 2 SO 4 BaSO KCl barium potassium barium potassium chloride sulfate sulfate chloride In order to determine which one of the products will be the precipitate requires a knowledge of the solubilities of salts in water. The rules governing the solubility of common salts are given below: THE SOLUBILITY RULES 1. All sodium, potassium, and ammonium salts are soluble in water. 2. The nitrates, chlorates, and acetates of all metals are soluble in water. Silver acetate is sparingly soluble. 3. The chlorides, bromides, and iodides of all metals except lead, silver, and mercury(i) are soluble in water. PbCl 2, PbBr 2, and PbI 2 are soluble in hot water. 6

7 4. The sulfates of all metals except lead, mercury(i), barium, and calcium are soluble in water. Ag 2 SO 4 is slightly soluble. 5. The carbonates, phosphates, borates, sulfites, chromates, and arsenates of all metal except sodium, potassium, and ammonium are insoluble in water. 6. The sulfides of all metals except barium, calcium, magnesium, sodium, potassium, and ammonium are insoluble in water. 7. The hydroxides of sodium, potassium, and ammonium are very soluble in water. The hydroxides of calcium and barium are moderately soluble. The oxides and hydroxides of all other metals are insoluble. b. Neutralization Reactions (sometimes called acid-base reactions) A neutralization reaction occurs between an acidic compound and a basic compound to form a chemical salt and water. 1. Reaction between an acid and a base HCl + NaOH NaCl + H 2 O hydrochloric sodium sodium water acid hydroxide chloride H 2 SO 4 + Mg(OH) 2 MgSO 4 + H 2 O sulfuric magnesium magnesium water acid hydroxide sulfate 2. Reaction between a metal oxide and an acid. When oxides of many metals are added to water, bases are formed. CaO + H 2 O Ca(OH) 2 calcium calcium oxide hydroxide (a metal oxide) (a base) Generally, these metal oxides are called basic anhydrides and they act like bases when mixed with acids. CaO + 2 HCl CaCl 2 + H 2 O calcium hydrochloric calcium oxide acid chloride Na 2 O + 2 HNO 3 2 NaNO 3 + H 2 O sodium nitric sodium oxide acid nitrate 3. Reaction between a non-metal oxide and a base. Many non-metal oxides are classified as acid anhydrides. These form acids when mixed with water. 7

8 SO 2 + H 2 O H 2 SO 3 sulfur sulfurous dioxide acid (a non-metal oxide) Non-metal oxides act as acids when mixed with a base. SO NaOH Na 2 SO 3 + H 2 O sulfur sodium sodium dioxide hydroxide sulfite CO 2 + Ca(OH) 2 CaCO 3 + H 2 O carbon calcium calcium dioxide hydroxide carbonate 5. Combustion Reactions Combustion reactions generally apply to organic compounds, such as hydrocarbons, which are used as fuels. In these cases, the compound is being burned in air (or oxygen) and producing carbon dioxide and water as products. A general form for a combustion reaction is: C n H 2n+2 + ( 3n + 1) O 2 n CO 2 + (n+1) H 2 O 2 Note: The actual coefficients will vary based on the composition of the starting compound. Some examples of combustion reactions are: C 3 H O 2 3 CO H 2 O propane 2 C 4 H O 2 8 CO H 2 O butane C 2 H 5 OH + 3 O 2 2 CO H 2 O ethanol 8

9 IV. THE ELECTROMOTIVE (ACTIVITY) SERIES OF METALS The Activity Series of Metals 1. Li - Lithium 2. K - Potassium 3. Ba - Barium 4. Sr - Strontium 5. Ca - Calcium 6. Na - Sodium 7. Mg - Magnesium 8. Al - Aluminum 9. Mn - Manganese 10. Zn - Zinc 11. Cr - Chromium 12. Fe - Iron 13. Cd - Cadmium 14. Co - Cobalt 15. Ni - Nickel 16. Sn - Tin 17. Pb - Lead 18. H - HYDROGEN 19. Sb - Antimony 20. As - Arsenic 21. Bi - Bismuth 22. Cu - Copper 23. Hg - Mercury(I) 24. Ag - Silver 25. Pd - Palladium 26. Hg - Mercury(II) 27. Pt - Platinum 28. Au - Gold Facts About the Activity Series a. The metals are arranged in the order of decreasing activity (i.e. their ability to pass into ionic form by losing electrons). Example: Lithium is more active than potassium, while potassium is more active than barium, etc... b. Each metal displaces any metal below it from dilute water solutions. Example: Fe + CuSO 4 FeSO 4 + Cu (Fe is above Cu in the activity series.) Fe + AlCl 3 No Reaction (Fe is below Al in the activity series.) c. Metals 1-6 react with cold water to liberate hydrogen, forming metal hydroxides. Metal 7 (Mg) displaces hydrogen from hot water. Example: Sr + 2 H 2 O Sr(OH) 2 + H 2 cold Co + H 2 O No Reaction cold d. Metals 1-13 react with steam to liberate hydrogen Example: 2 Al + 3 H 2 O Al 2 O H 2 steam Ni + H 2 O No Reaction steam e. Metals 1-17 react with acids to liberate hydrogen Example: Sn + 2 HCl SnCl 2 + H 2 Hg + H 2 SO 4 No Reaction 9

10 Reaction of metals with nitric acid results in a decomposition of the nitric acid along with a displacement reaction. With active metals such as iron and zinc, the reaction with concentrated nitric acid is: 4 Zn + 10 HNO 3 4 Zn(NO 3 ) 2 + N 2 O + 5 H 2 O With less active metals, such as copper, the reaction with concentrated nitric acid is: Cu + 4 HNO 3 Cu(NO 3 ) NO H 2 O Concentrated nitric acid will react with metals similar to the reaction with copper, above. f. Metals 1-23 react with oxygen to form oxides. The oxides of Ag, Pd, Pt, and Au can be prepared only by indirect methods. Example: 2 Fe + O 2 Cu + H 2 O Ag + O 2 No Reaction g. The oxides of metals can be reduced by hydrogen to yield the metal and water. The other oxides cannot be reduced by hydrogen. Example: CuO + H 2 Cu + H 2 O MgO + H 2 No Reaction h. The oxides of metals can be decomposed by the heat of a Bunsen burner. The other oxides cannot be decomposed by the heat of a Bunsen burner. Example: 2 HgO + heat 2 Hg + O 2 Al 2 O 3 + heat No Reaction i. It is to be noted that the most active elements form the most stable compounds. Example: 2 HgO + heat 2 Hg + O 2 ZnO + heat No Reaction Since Zn is more active than Hg, ZnO is more stable than HgO and it cannot be decomposed by simple heating. NOTE: Examine the list of elements in the Activity Series on the preceding page. Rather than trying to memorize the entire list, the order of the elements in the series can be related to the periodic table. In general, Group IA elements are at the top followed by Group IIA elements and then Group IIIA elements. Next are the common transition elements, then Group IVA elements followed by hydrogen. Below hydrogen are the Group VA elements followed by elements used for dental fillings and jewelery (Group IB and neighbors ). 10

11 V. THE EFFECT OF HEAT ON METALLIC COMPOUNDS: PREDICTION OF PRODUCTS OF DECOMPOSITION REACTIONS 1. On Oxides The oxides of Fe, Cd, Co, Ni, Sn Pb, Sb, As, Bi, Cu, Hg, Ag, Pd, Pt, and Au (metals on the electromotive series) can be reduced by hydrogen to yield the metal and water. The other oxides cannot be reduced by hydrogen. Example: Fe 2 O H 2 2 Fe + 3 H 2 O The oxides of Hg, Ag, Pd, Hg, Pt, and Au (metals on the electromotive series) can be decomposed by the heat of a Bunsen burner. The other oxides cannot be decomposed by the heat of a Bunsen burner. Example: 2 Ag 2 O + heat 4 Ag + O 2 2. On Hydroxides All hydroxides, except those of the alkali metals (Group IA), will lose water when heated forming the metal oxide. Examples: Mg(OH) 2 + heat MgO + H 2 O NaOH + heat No Reaction The hydroxides of mercury and silver are not stable, they decompose to form the oxide and water without heating. Examples: AgNO 3 + NaOH AgOH + NaNO 3 3. On Sulfates 2 AgOH Ag 2 O + H 2 O With the exception of the alkali metal (Group IA) and alkaline earth (Group IIA) sulfates, the sulfates of all other metals are decomposed by heat to form the metal oxide and sulfur trioxide. Examples: Al 2 (SO 4 ) 3 + heat Al 2 O SO 3 4. On Nitrates CaSO 4 + heat No Reaction The nitrates of the alkali metals decompose on heating to yield the nitrites and oxygen. All other metal nitrates are decomposed to nitrogen dioxide, oxygen, and the metal oxide on heating. Examples: 2 KNO 3 + heat 2 KNO 2 + O 2 2 Pb(NO 3 ) 2 + heat 2 PbO + 4 NO 2 + O 2 11

12 5. On Carbonates Except for the alkali metal carbonates, all carbonates lose carbon dioxide when heat to form the metal oxide. Examples: MgCO 3 + heat MgO + CO 2 Na 2 CO 3 + heat No Reaction NOTE ON SECTIONS 2-5: The hydroxides, sulfates, nitrates, and carbonates of metals in the activity series will yield the metal on heating, since the oxides of these metals are decomposed by heat. For example, the reaction of Au(OH) 3 will also cause the decomposition of the Au 2 O 3 : 2 Au 2 O 3 + heat 4 Au + 3 O 2 The overall reaction, combining the above two steps, can be written: 4 Au(OH) 3 + heat 4 Au + 6 H 2 O + 3 O 2 As another example, consider the decomposition of silver sulfate on heating. The overall reaction is: 2 Ag 2 SO 4 + heat 4 Ag + 2 SO 3 + O 2 6. On Chlorates All chlorates decompose on heating to form the chloride of the metal and oxygen gas. 2 KClO 3 + heat 2 KCl + O 2 Thermal decomposition of the bromates and iodates result in a number of different products depending on the conditions under which the reactions occur. No general rule can be written for the decomposition of these compounds due to heating. 12

13 SAMPLE PROBLEMS Complete and balance the following equations. If no reaction takes place, indicate by N.R. 1. ZnSO 4 + heat Solution: This reaction should be recognized as a decomposition reaction (a single compound plus heat). Looking at Section 12.5, concerning the action of heat of sulfates (page 92), it is found that ZnSO 4 should decompose to form the oxide and SO 3. The completed equation should be: This equation is balanced as written. ZnSO 4 + heat ZnO + SO 3 2. Sn + CdCl 2 Solution: A reaction occurring between an element and a compound fits the form of a displacement reaction. Referring to the activity series on page 90, it is observed that Sn is below Cd. Thus Sn is not active enough to replace Cd and no reaction will take place. Sn + CdCl 2 N.R. 3. ZnO + H 3 PO 4 Solution: This reaction is occurring between a metal oxide (a basic anhydride) and an acid. Therefore, this is a neutralization reaction or a form of the metathesis type reaction. The completed reaction will be: ZnO + H 3 PO 4 Zn 3 (PO 4 ) 2 + H 2 O The completed equation must be balanced. (see pages 83-84) The final balanced equation will be: 3 ZnO + 2 H 3 PO 4 Zn 3 (PO 4 ) H 2 O 13

14 PROBLEMS: Writing chemical equations. 1. Complete and balance the following direct union equations. a) K + Br 2 b) Mg + O 2 c) H 2 + Br 2 d) Na + I 2 e) CaO + SO 2 f) Zn + O 2 g) Na 2 O + SO 3 h) N 2 + H 2 i) Cu + S j) H 2 O + P 2 O 5 2. Complete and balance the following decomposition equations. If no reaction takes place, indicate by writing N.R. a) NaNO 3 + heat b) CaCO 3 + heat c) HgSO 4 + heat d) NaOH + heat e) Hg(NO 3 ) 2 + heat f) KClO 3 + heat g) PbSO 4 + heat h) Fe(OH) 3 + heat i) Ag 2 CO 3 + heat j) Ba(ClO 3 ) 2 + heat 14

15 3. Complete and balance the following displacement equations. If no reaction takes place, indicate by writing N.R. a) Zn + H 2 SO 4 b) Cr + PbCl 2 c) Ag + HCl d) Al + CuSO 4 e) Li + H 2 O f) Cl 2 + KBr g) Ni + H 2 O (steam) h) Cu + H 2 SO 4 i) Pb + FeCl 3 j) Zn + SnBr 2 4. Complete and balance the following metathesis equations. a) Al 2 O 3 + HNO 3 b) HgNO 3 + HCl c) NiSO 4 + Na 2 CO 3 d) Pb(NO 3 ) 2 + HCl e) Cr 2 O 3 + H 2 SO 4 f) Ca(OH) 2 + H 3 PO 4 g) CuSO 4 + H 2 S h) FeCl 3 + Ca(OH) 2 i) AgNO 3 + Na 2 CrO 4 j) Al(OH) 3 + HCl 15

16 5. Complete and balance the following combustion equations. a) CH 4 + O 2 b) C 5 H 12 + O 2 c) C 8 H 18 + O 2 d) C 3 H 7 OH + O 2 e) C 2 H 5 OC 2 H 5 + O 2 6. Classify each of the following equations as direct union, decomposition, displacement, or metathesis reactions AND complete and balance each equation. If no reaction takes place, indicate by N.R. a) CuSO 4 + heat type: b) CrCl 3 + Na 2 SiO 3 type: c) Fe 2 O 3 + H 2 type: d) MgO + CO 2 type: e) Ag + H 2 SO 4 type: f) NaOH + HNO 3 type: g) K 2 SO 4 + heat type: h) Ni + Pt(SO 4 ) 2 type: i) Fe + H 2 O (steam) type: j) Ba + F 2 type: k) KOH + CO 2 type: l) Pt + O 2 type: m) Ba(OH) 2 + H 2 CO 3 type: n) Ni(NO 3 ) 2 + heat type: o) Sr + H 2 O type: p) Ag 2 CO 3 + heat type: q) MnCl 2 + H 2 S type: r) CaO + HNO 3 type: s) Cu + HCl type: 16

17 t) Al + Br 2 type: u) Pb(NO 3 ) 2 + Na 2 CrO 4 type: v) Al(OH) 3 + heat type: w) Ca(OH) 2 + H 3 PO 4 type: x) AsCl 3 + H 2 S type: y) Cu + Hg(NO 3 ) 2 type: z) Co 2 (SO 4 ) 3 + heat type: aa) Fe + HNO 3 bb) Ni(NO 3 ) 2 + NaOH cc) Na + H 2 O dd) KCl + H 2 SO 4 ee) Mg + H 2 O (steam) type: type: type: type: type: ff) Ag + HNO 3 type: gg) CaCO 3 + HCl hh) KOH + H 2 SO 4 type: type: ii) SrI 2 + Br 2 type: jj) Ba(NO 3 ) 2 + heat type: kk) Al + H 2 SO 4 type: ll) PCl 3 + Cl 2 type: mm) SO 2 + O 2 type: oo) K 3 PO 4 + BaCl 2 type: 17